Types of Chemical Reactions (rxns.) - PowerPoint PPT Presentation

types of chemical reactions rxns n.
Download
Skip this Video
Loading SlideShow in 5 Seconds..
Types of Chemical Reactions (rxns.) PowerPoint Presentation
Download Presentation
Types of Chemical Reactions (rxns.)

play fullscreen
1 / 45
Types of Chemical Reactions (rxns.)
198 Views
Download Presentation
eadoin
Download Presentation

Types of Chemical Reactions (rxns.)

- - - - - - - - - - - - - - - - - - - - - - - - - - - E N D - - - - - - - - - - - - - - - - - - - - - - - - - - -
Presentation Transcript

  1. 0 Types of Chemical Reactions (rxns.)

  2. Introduction 0 • Chemical reactions occur when bonds (between the electrons of atoms) are formed or broken • Chemical reactions involve • changes in the chemical composition of matter (the making of new materials with new properties) • energy changes • Symbols represent elements • Formulas describe compounds • Chemical equations describe a chemical reaction

  3. Types of Reactions 0 • Reactions are classified by their products. • There are five types of chemical reactions we will talk about: • Synthesis reactions • Decomposition reactions • Single displacement reactions • Double displacement reactions • Combustion reactions • You need to be able to identify the type of reaction and predict the product(s)

  4. Steps to Writing Reactions 0 Some steps for doing reactions: • Identify the type of reaction • Predict the product(s) using the type of reaction as a model • Balance it Don’t forget about the diatomic elements! (BrINClHOF) For example, Oxygen is O2 as anelement. In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s a compound!

  5. Synthesis Reactions 0 • Synthesis (meaning to make) are • Also called • Direct combination or combination reactions • Addition reactions • typified by their single product. • If you have a reaction in which at least 2 elements or compounds are reacted and produce a single product, the reaction is a synthesis reaction.

  6. Synthesis Reactions 0 reactant + reactant  1 product • Basically: A + B  AB • Example: 2H2 + O2  2H2O • Example: C+ O2  CO2 • Note: Single Product! This is your clue that this is a synthesis or combination reaction.

  7. Synthesis Reactions 0 • Here is another example of a synthesis reaction

  8. Examples of Synthesis Reactions • 2Na + S Na2S • This one is an example of two elements in atomic form (Na and S) combining to form a compound (sodium sulfide). • 2H2 + O2 2H2O • In this example, A and B are two elements in molecular form (hydrogen and oxygen molecules), and the product is water, which is simply the chemical combination of hydrogen and oxygen.

  9. Examples of Synthesis Reactions • 4Fe + 3O2 2Fe2O3 • In this example, substance “A” is an element in atomic form (Fe), and substance “B” is an element in molecular form (O2). The result is a direct chemical combination of the two elements (FeO, iron oxide, which is “rust”). • CuO + H2O Cu(OH)2 • This is an example where both substances going into the reaction are molecules. The result is what you get when you “add” all of the atoms in the reaction together.

  10. Practice 0 • Predict the products. Write and balance the following synthesis reaction equations. • Sodium metal reacts with chlorine gas Na(s) + Cl2(g)  • Solid Magnesium reacts with fluorine gas Mg(s) + F2(g)  • Aluminum metal reacts with fluorine gas Al(s) + F2(g)  2 2 NaCl(s) MgF2(s) Balanced 2 3 2 AlF3(s)

  11. Decomposition Reactions 0 Decomposition reactions are really just the opposite of a synthesis reaction. Remember, if you can make a substance, you should be able to break it back apart into its components. A good way to remember decomposition reactions to to remember what happens when something decomposes. It falls apart!

  12. Decomposition Reactions 0 • Decomposition reactions occur when a compound breaks up into the elements or in a few to simpler compounds • 1 Reactant  Product + Product • Basically: AB  A + B • Example: 2 H2O  2H2 + O2 • Example: 2 HgO  2Hg + O2 • Note: Single Reactant! The single reactant is your clue that this is a decomposition reaction.

  13. Decomposition Reactions 0 • Another view of a decomposition reaction:

  14. Decomposition Exceptions 0 • Carbonates and chlorates are special case decomposition reactions that do not go to the elements. • Carbonates (CO32-) decompose to carbon dioxide and a metal oxide • Example: CaCO3  CO2 + CaO • Chlorates (ClO3-) decompose to oxygen gas and a metal chloride • Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2 • There are other special cases, but we will not explore those in this class

  15. Practice 0 • Predict the products. Then, write and balance the following decomposition reaction equations: • Solid Lead (IV) oxide decomposes PbO2(s)  • Aluminum nitride decomposes AlN(s)  Pb(s) + O2(g) 2 2 Al(s) + N2(g)

  16. Practice 0 Identify the type of reaction for each of the following synthesis or decomposition reactions, and write the balanced equation: N2(g) + O2(g) BaCO3(s)  Co(s)+ S(s)  NH3(g) + H2CO3(aq)  2NI3(s)  2 NO (g) BaO(s) + CO2 (g) CoS (s) (NH4)2CO3(s) N2 (g) + 3I2 (s)

  17. Single Replacement Reactions 0 Single replacement reactions occur when one chemical takes the place of another in a reaction. In the typical single replacement reaction, an element trades places with one of the ions in a compound.

  18. Single Replacement Reactions 0 • A metal can replace a metal (+) OR a nonmetal can replace a nonmetal (-). • element + compound product + product A + BC  AC + B (if A is a metal)OR A + BC  BA + C (if A is a nonmetal) (remember the cation always goes first!) When H2O splits into ions, it splits into H+ and OH- (not H+ and O-2 !!)

  19. Single Replacement Reactions 0 • Another view:

  20. The Activity Series 0 • Not all single replacement reactions will occur. • This depends upon the location of the elements present in the activity series • Elements above MAY replace elements below; elements below MAY NOT replace elements above them on the series

  21. 0 You will be given a copy of this!!!!

  22. Single Replacement Reactions 0 • Write and balance the following single replacement reaction equation: • Zinc metal reacts with aqueous hydrochloric acid Zn(s) + HCl(aq) ZnCl2 + H2(g) Note: Zinc replaces the hydrogen ion in the reaction [If ZnCl2 + H2(g)  Zn(s) + HCl(aq) the reaction WOULD NOT OCCUR because Hydrogen is below zinc on the activity series] 2

  23. Single Replacement Reactions 0 2 2 • Sodium chloride solid reacts with fluorine gas NaCl(s) + F2(g)  NaF(s) + Cl2(g) Note that fluorine replaces chlorine in the compound • Aluminum metal reacts with aqueous copper (II) nitrate Al(s)+ Cu(NO3)2(aq) Cu(s) + Al(NO3)3(aq) 2 3 3 2

  24. Double Replacement Reactions 0 Double replacement reactions are identified by two ions trading places and forming new compounds.

  25. Double Replacement Reactions 0 • Double Replacement Reactions occur when a metal replaces a metal in a compound and a nonmetal replaces a nonmetal in a compound • Compound + compound  product + product • AB + CD  AD + CB • Notice that one ion from compound AB replaces one ion from compound CD.

  26. Double Replacement Reactions 0 • Think about it like “foil”ing in algebra, first and last ions go together + inside ions go together • Example: AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq) • Another example: K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

  27. Solubility 0 • For a double replacement reaction to have occurred, a solid (precipitate) MUST be formed • There are rules to determine which of the materials formed is the solid • If no solid is formed, there is said to be no reaction.

  28. Figure 8.4: The forming of solid AgCl. 0

  29. Solubility Tables 0 • Solubility tables help determine which materials are soluble in water and which are not • In general, Solubility Rules can be summarized as follows 1. All compounds containing alkali metal cations and the ammonium ion are soluble. 2. All compounds containing NO3-, ClO4-, ClO3-, and C2H3O2- anions are soluble. 3. All chlorides, bromides, and iodides are soluble except those containing Ag+, Pb2+, or Hg22+. 4. All sulfates are soluble except those containing Hg22+, Pb2+, Sr2+, Ca2+, or Ba2+. 5. All hydroxides are insoluble except compounds of the alkali metals, Ca2+, Sr2+, and Ba2+. 6. All compounds containing PO43-, S2-, CO32-, and SO32- ions are insoluble except those that also contain alkali metals or NH4+. You will be given a copy of this!!!!

  30. Practice 0 • Predict the products. Balance the equation • HCl(aq) + AgNO3(aq)  • Pb(NO3)2(aq) + BaCl2(aq)  • FeCl3(aq) + 3NaOH(aq)  • H2SO4(aq) + 2NaOH(aq)  HNO3(aq) + AgCl(s) PbCl2(s) + Ba(NO3)2(aq) Fe(OH)3(s) + 3NaCl(aq) 2 H2O(l) + Na2SO4(aq) NR

  31. Combustion Reactions 0 Combustion reactions are the ones that burn (or explode!). There are two types of combustion reactions—complete or incomplete reactions. These reactions are identified by their products. They either produce carbon monoxide and water or carbon dioxide and water.

  32. Complete Combustion Reactions 0 These reactions burn “efficiently” which means they produce carbon dioxide and water. These reactions typically burn cleanly and leave very little residue behind.

  33. Combustion Reactions 0 • Combustion reactions occur when a hydrocarbon reacts with oxygen gas. • This is also called burning!!! In order to burn something you need the 3 things in the “fire triangle”:1) A Fuel (hydrocarbon)2) Oxygen to burn it with3) Something to ignite the reaction (spark)

  34. Combustion Reactions 0 • In general: CxHy + O2  CO2 + H2O • Products in combustion are ALWAYS carbon dioxide and water. (although incomplete burning does cause some by-products like carbon monoxide) • Combustion is used to heat homes and run automobiles (octane, as in gasoline, is C8H18)

  35. Complete Combustion Reactions 0 • CH4 + O2 CO2 + H2O • They may also be written: CH4  CO2 + H2O With O2 written above the arrow • Clues: CO2 (carbon dioxide) in the product along with water

  36. Combustion 0 • Example • C5H12 + O2 CO2 + H2O • Write the products and balance the following combustion reaction: • C10H22 + O2  • C10H22 + O2  8 5 6 15.5 10 CO2 + H2O 11 2 31 20 CO2 + H2O 22

  37. Incomplete Combustion Reactions 0 Incomplete combustion reactions occur when something does not burn efficiently. This can cause a lot of harm if the gases produced cannot escape. Carbon monoxide,an odorless and colorless gas, is dangerous. People poisoned by this gas usually become sleepy and can die due to exposure.

  38. Incomplete Combustion Reactions 0 • CH4 + O2 CO + H2O • These reactions may also be written by: CH4 CO + H2O Again, the O2 is usually written over the arrow. • Clue: CO (Carbon monoxide as a product.)

  39. Combustion Reactions 0 Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning.

  40. Mixed Practice 0 • State the type, predict the products, and balance the following reactions: • BaCl2 + H2SO4 • C6H12 + O2  • Zn + CuSO4  • Cs + Br2  • FeCO3  BaSO4 + HCl CO2 + H2O ZnSO4 + Cu CsBr FeO + CO2

  41. Predicting Products of Reactions 0 • Completing reactions requires knowledge of the different reaction types (sometimes called mechanisms). • You must first identify the reaction type by the reactants. • The only type of reaction that cannot be predicted this way is the combustion reaction since the products are very similar.

  42. First Step: 0 • Identify reaction type Example: Al + O2  Clue: 2 elements – Synthesis or combination reaction

  43. Second Step: 0 • Write the net ionic equation for the reactants Al + O2 becomes Al3+ + O2- 

  44. Step 3 0 • Using clues, complete reaction taking care to write each formula correctly by checking charges and “criss-crossing” if necessary. Al + O2 Al3+O2- Al + O2 Al2O3

  45. Predicting Products of Reactions (cont.) 0 • For Single Replacement reactions, check activity series to make sure the reaction goes. • Once you write the molecular equation, you should check for reactants and products that are soluble or insoluble. (Double Replacement only)